Study Of The Reaction Between Iron (III) And Thiocyanate Ions

Reaction between Iron (III) ( and Thiocyanate ions (.

Discussion and Conclusion

The present’s set-up was to study the reaction between Iron (III) ( and thiocyanate ions (, in an aqueous solution that reduces a complex ion,  via a ligand exchange reaction. Through quantitative preparations of various solutions and successive measurement of the absorbance of the solution utilizing a spectrometer, the concentration of was obtained. The difference between the initial concentration and the used/consumed concentration by the generation of  enabled for calculation of  (concentrations. The equilibrium constant for the formation of the complex was calculated from the combined concentration

Calibration curve was prepared by running a series of solutions with known concentrations and the absorbance readings were taken, and then plotting the absorbance reading as a function of thiocyanate concentration. The calibration curve yielded a straight line originating from the origin.  By applying the beer’s lamber law, the concentration of  was evaluated (using equation of the calibration curve. The concentration of  at equilibrium was calculated by subtracting the  from the initial concentration of Fe. The equilibrium concentration of  was calculated by subtracting the  from the initial concentration of thiocyanate. The equilibrium constant for the reaction was then evaluated as;

This was done for every prepared liquot solution. The equilibrium constant values were then averaged.

By literature, the value of  in relation to 1 (that is, < or >) signifies some chemical properties. Being a ratio, it can tell is whether there are less reactants or products at equilibrium for a particular reaction.  When the experimentally obtained  value is greater than 1, then it implies that the concentration of the products is significantly more than that of the reactants. That is, the reaction goes to completion; if not all, most of the reactants are consumed up to generate products. On the other hand, if the equilibrium constant is less 1, then it implies that the concentration of products is much less and therefore, the reaction doesn’t proceed to greater extent; less products are formed. The experimentally calculated  is greater than 1 which indicates that more were formed.

The significant figures in the equilibrium constant were greater than the significant figures in the measurements implying that the uncertainty of the of the established equilibrium constant is greater than the uncertainty in the measurements, indicating the presence of an extra source of random errors in the uncertainty of the obtained results. Another possible source of error (systematic error) could be on the concentration of the stock KSCN reagent and the spectrophotometer calibration.

Conclusion; the equilibrium constant of was successfully determined. The obtained constant delinquently corroborates and support the specific target of the experiment. Though there were some errors that were associated with the set-up, the observed value forms a strong basis that support the experimental aim.